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Fluoroantimonic acid is an inorganic compound with the chemical formula H
(also written H
, 2HF·SbF5, or simply HF-SbF5). It is an extremely strong acid, easily qualifying as a superacid. The Hammett acidity function, H0, has been measured for different ratios of HF:SbF5. While the H0 of pure HF is −15, addition of just 1 mol % of SbF5 lowers it to around −20. However, further addition of SbF5 results in rapidly diminishing returns, with the H0 reaching −21 at 10 mol %. The use of an extremely weak base as indicator shows that the lowest attainable H0, even with > 50 mol % SbF5, is somewhere between −21 and −23.[1][2][3]

Fluoroantimonic acid
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Fluoroantimonic acid 3D spacefill.png
IUPAC name
Fluoroantimonic acid
Systematic IUPAC name
Fluoranium hexafluorostibanuide
Fluoranium hexafluoridoantimonate(1−)
3D model (JSmol)
ECHA InfoCard 100.037.279
EC Number 241-023-8
Molar mass 256.765


Appearance Colorless liquid
Density 2.885 g/cm3
Solubility SO2ClF, SO2
Main hazards Extremely corrosive, Violent hydrolysis
GHS pictograms GHS05: CorrosiveGHS06: ToxicGHS07: HarmfulGHS09: Environmental hazard
GHS signal word Danger
H300, H310, H314, H330, H411
P260, P264, P273, P280, P284, P301+310
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 4: Very short exposure could cause death or major residual injury. E.g., VX gasReactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g., fluorineSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
Related compounds
Related acids
Antimony pentafluoride

Hydrogen fluoride
Magic acid

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

The "canonical" composition of fluoroantimonic acid is prepared by treating hydrogen fluoride (HF) with antimony pentafluoride (SbF5) in a stoichiometric ratio of 2:1. It is the strongest superacid based on measured H0 value. Only the carborane acids, whose H0 could not be directly determined due to their high melting points, may be stronger acids than fluoroantimonic acid.[4] It has been shown to protonate even hydrocarbons to afford pentacoordinate carbocations (carbonium ions).[5][6]

The reaction to produce fluoroantimonic acid results in formation of the fluoronium ion:

SbF5 + 2 HF → SbF
+ H2F+

The acid is often said to contain "naked protons", but the "free" protons are, in fact, always bonded to hydrogen fluoride molecules.[7] It is the fluoronium ion that accounts for fluoroantimonic acid's extreme acidity. The protons easily migrate through the solution, moving from H2F+ to HF, when present, by the Grotthuss mechanism:

H2F+ + HF ⇌ HF + H2F+

Fluoroantimonic acid thermally decomposes at higher temperatures, generating free hydrogen fluoride gas. It is exceptionally corrosive and can only be stored in containers lined with Teflon.



Two related products have been crystallized from HF-SbF5 mixtures, and both have been analyzed by single crystal X-ray crystallography. These salts have the formulas [H
and [H
. In both salts, the anion is Sb
.[8] As mentioned above, SbF
is weakly basic; the larger anion Sb
is expected to be still weaker.

The following values show that fluoroantimonic acid is much stronger than other superacids[4] based upon the Hammett acidity function. Increased acidity is indicated by smaller (in this case, more negative) values of H0.


This extraordinarily strong acid deprotonates nearly all organic compounds. In 1967, Bickel and Hogeveen showed that 2HF·SbF5 will remove H2 from isobutane and methane from neopentane to form carbenium ions:[9][10]

(CH3)3CH + H+ → (CH3)3C+ + H2
(CH3)4C + H+ → (CH3)3C+ + CH4

It is also used in the manufacture of tetraxenon gold compounds.

Materials compatible with fluoroantimonic acid as a solvent include SO2ClF, and sulfur dioxide; some chlorofluorocarbons have also been used. Containers for HF-SbF5 are made of PTFE.


HF-SbF5 is extremely corrosive, toxic, and moisture sensitive.[6] Like most strong acids, fluoroantimonic acid can react violently with water, owing to the exothermic hydration. Consequently, it cannot be used in aqueous solution, only in hydrogen fluoride as solvent.

Fluoroantimonic acid cannot be stored in glass, as it will dissolve it. It must be stored in a PTFE (polytetrafluoroethylene) container.[11]

See alsoEdit


  1. ^ Superacid chemistry. Olah, George A. (George Andrew), 1927-2017., Olah, George A. (George Andrew), 1927-2017. (2nd ed.). Hoboken, N.J.: Wiley. 2009. ISBN 9780470421543. OCLC 391334955.CS1 maint: others (link)
  2. ^ Olah, G. A. (2005). "Crossing Conventional Boundaries in Half a Century of Research". Journal of Organic Chemistry. 70 (7): 2413–2429. doi:10.1021/jo040285o. PMID 15787527.
  3. ^ In ref. 2 (2005), Olah estimates that HF-SbF5 may reach H0 values as low as –28. On the other hand, in ref. 1 (2009), Olah cites one method that estimated H0 values down to –27 for FSO3H-SbF5 at 90% SbF5, but indicates that more reliable experimentally determined equilibrium constants do not support H0 values lower than about –24 for either magic acid or fluoroantimonic acid.
  4. ^ a b Gillespie, R. J.; Peel, T. E. (1973-08-01). "Hammett acidity function for some superacid systems. II. Systems sulfuric acid-[fsa], potassium fluorosulfate-[fsa], [fsa]-sulfur trioxide, [fsa]-arsenic pentafluoride, [sfa]-antimony pentafluoride and [fsa]-antimony pentafluoride-sulfur trioxide". Journal of the American Chemical Society. 95 (16): 5173–5178. doi:10.1021/ja00797a013. ISSN 0002-7863.
  5. ^ Olah, G. A. (2001). A Life of Magic Chemistry: Autobiographical Reflections of a Nobel Prize Winner. John Wiley and Sons. pp. 100–101. ISBN 978-0-471-15743-4.
  6. ^ a b Olah, G. A.; Prakash, G. K. Surya; Wang, Qi; Li, Xing-ya (15 April 2001). "Hydrogen Fluoride-Antimony(V) Fluoride". Hydrogen Fluoride–Antimony(V) Fluoride. Encyclopedia of Reagents for Organic Synthesis. New York: John Wiley and Sons. doi:10.1002/047084289X.rh037m. ISBN 9780470842898.
  7. ^ Klein, Michael L. (October 25, 2000). "Getting the Jump on Superacids" (PDF). Pittsburgh Supercomputing Center (PSC). Retrieved 2012-04-15.
  8. ^ Mootz, Dietrich; Bartmann, Klemens (March 1988). "The Fluoronium Ions H2F+ and H
    : Characterization by Crystal Structure Analysis". Angewandte Chemie International Edition. 27 (3): 391–392. doi:10.1002/anie.198803911.
  9. ^ Bickel, A. F.; Gaasbeek, C. J.; Hogeveen, H.; Oelderik, J. M.; Platteeuw, J. C. (1967). "Chemistry and spectroscopy in strongly acidic solutions: reversible reaction between aliphatic carbonium ions and hydrogen". Chemical Communications. 1967 (13): 634–635. doi:10.1039/C19670000634.
  10. ^ Hogeveen, H.; Bickel, A. F. (1967). "Chemistry and spectroscopy in strongly acidic solutions: electrophilic substitution at alkane-carbon by protons". Chemical Communications. 1967 (13): 635–636. doi:10.1039/C19670000635.
  11. ^ "What Is the World's Strongest Superacid?". ThoughtCo. Retrieved 2017-04-06.